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WATER
COURSE OUTCOME (C0 1)
 CO2: Ability to define and describe the




biochemical concepts and terms associated
with life.
Terms used in Course Outcome and
Teaching
Knowledge: Define, introduce, describe,
name, relate, explain, identify and Remember
concepts and principles.
Repetition: Repeat and discuss concepts and
principles.
Application: Apply, demonstrate, interpret and
illustrate concepts and principles.
*LECTURE CONTENTS
Sect
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
13.
14.
Why water is important to biochem ?
USES OF WATER
PHYSICS & CHEMISTRY OF WATER
PHYSICAL PROPERTIES OF WATER
Molecular Structure of Water
Hydrogen Bonding
Polarity of water
Noncovalent Bonding
Weak, van der Waal’s forces
Thermal properties of water
Osmosis, reverse osmosis & dialysis
Water ionization, pH, titration and buffer
Summary
The end note
*Sect 1. Introduction
Why water is important to biochem ?

More than 70% of the earth’s surface is
covered with the molecule of water.
 Cell components and molecules (protein,
poly sacharides, nucleic acid, membranes)
assume their shape in response to water
 Water acts as a solvent & substrate for
many cellular reactions

Water is a common chemical substance that is
essential for the survival of all known forms of
life. (In typical usage, water refers only to its liquid form or state, but the
substance also has a solid state, ice, and a gaseous state, water vapor.
)
Introduction
Uses of Water
 AGRICULTURE
 FOR DRINKING
 AS SOLVENT
 HEAT TRANSFER FLUID
 FOOD PROCESSING
 INDUSTRIAL APPLICATIONS
 AS A SCIENTIFIC STANDARD
Introduction
Physics & Chemistry of water
 Water is the chemical substance with chemical formula
H2O: one molecule of water has two hydrogen atoms
covalently bonded to a single oxygen atom.
 Water is a tasteless, odorless liquid at ambient
temperature and pressure, and appears colorless in small
quantities, although it has its own intrinsic very light blue
hue.
 Oxygen attracts electrons much more strongly than
hydrogen, resulting in a net positive charge on the
hydrogen atoms, and a net negative charge on the
oxygen atom.
 The presence of a charge on each of these atoms gives
each water molecule a net dipole moment.
 Electrical attraction (hydrogen bonding) between water
molecules due to this dipole pulls individual molecules
closer together, making it more difficult to separate the
molecules and therefore raising the boiling point.
 Water can be described as a polar liquid that dissociates
disproportionately into the hydronium ion (H3O+(aq)) and
an associated hydroxide ion (OH−(aq))
 Water is in dynamic equilibrium between the liquid, gas
and solid states at standard temperature and pressure
(0°C, 100.000 kPa) , and is the only pure substance
found naturally on Earth.
Introduction
UNIQUE PHYSICAL PROPERTIES OF WATER
 Exist in all three physical states of
matter: solid, liquid, and gas.
 Has high specific heat
 Water conducts more easily than any
liquid except mercury
 Water has a high surface tension
 Water is a universal solvent
 Water in a pure state has a neutral pH.
*Sect 2 : Molecular Structure of Water
 The oxygen in
water is sp3
hybridized.
Therefore water
has tetrahedral
geometry.
Consequently the
water molecule is
bent. The H-O-H
angle is 104.5o.
Molecular structure of water
 The bent structure indicate water is polar
coz linear structure is nonpolar.
 Phenomenon where charge is separated
to partial –ve charge and partial +ve
charge is called dipoles.
Sect 3 : Noncovalent Bonding
 Ionic interactions
 Hydrogen bonding
 Van der Waals forces
 Dipole-dipole
 Dipole-induced
dipole
 Induced dipole-induced dipole
Ionic Interactions
 Ionic interactions occur between charged atoms or
groups.
 Oppositely charged ions are attracted to each
other.
 In proteins, side chains sometimes form ionic salt
bridges.
Salt bridge
-
CH2CH2COO
+
H3N CH2CH2
Hydrogen Bonding
 Water molecules can perform hydrogen bond with
one another. Four hydrogen bonding attractions
are possible per molecule:
2 through the
hydrogens and
H
H
2 through the
O
O
nonbonding
H
H
electron pairs.
H
O
H
H
O
H
H
O
H
Hydrogen Bonding
 A hydrogen attached to an O
becomes very polarized and highly
partial plus (δ+). This partial
positive charge interacts with the
nonbonding electrons on another
O giving rise to the very powerful
hydrogen bond.
hydrogen bond
shown in yellow
O
H
H
R1 O H
H
O
H
 Water has an abnormally high boiling point
due to intermolecular hydrogen bonding.
H 

O
H


H 
O
H

H

O

H

H bonding is a
weak attraction
between an
electronegative
atom in one
molecule and an H
(on an O) in
another.
 Water molecule with bond
(
(
) and net
) dipoles.
-
H
+
O
H +
Van der Waal’s forces
 These forces are electrostatic interactions.
 Relatively weak.
 These interactions occur among permanent or
induced dipoles, where the hydogen bonds form a
special kind of dipolar interaction
 Interactions among permanent dipoles such as
carbonyl groups are much weaker than ionic
interactions
Van der Waals Attractions
a. Dipole-dipole
b. Dipole-induced dipole
c. Induced dipole-induced dipole
+ C
-
O
+ C
-
O
H
H
+ C
-
O
H
+
H
H
+
+
H -
H
H
H - H
H
H -
Sect 4 : Thermal properties of water
 Hydrogen bonding keeps water in the liquid phase
between 0oC and 100oC.
 Liquid water has a high:
Heat of vaporization - energy to vaporize one mole
of liquid at 1 atm
Heat capacity - energy to change the
temperature by 1oC
 Water plays an important role in thermal regulation in
living organisms.
Relationship between temperature and hydrogen
bond
 Max number of hydrogen bonds form when water has




frozen into ice.
Hydrogen bonds is approximately 15% break when ice
is warmed.
Liquid water consists of continuously breaking and
forming hydrogen bonds.
The rising tempt. The broken of hydrogen bonds are
accelerating.
When boiling point is reached, the water molecules
break free from one another and vaporize.
Sect. 5: Solvent Properties of Water
 Water easily dissolves a wide variety of
the constituents of living organisms.
 Water also unable to dissolve some
substances
 This behavior is called hydrophilic and
hydrophobic properties of water.
Polarity of water
 Water is a polar molecule.

A polar molecule is one in which one end is
partially positive and the other partially
negative.

This polarity results from unequal sharing of
electrons in the bonds and the specific
geometry of the molecule.
Hydrophillic substances
dissolve in water
 The polar nature of water makes it an excellent
solvent for polar and ionic materials that are
water loving (hydophillic)
 An ion immersed in a polar solvent such as
water attracts the the oppositely charged ends
of the solvent dipoles and becomes surrounded
by one or more concentric shells of orientated
solvent molecules, therby becoming solvated or
hydrated when water is the solvent
 The bond dipoles of uncharged polar molecules
make them soluble in aqueous solutions.
HYDROPHOBIC INTERACTIONS
 Nonpolar molecules tend to coalesce
into droplets in water. The repulsions
between the water molecules and the
nonpolar molecules cause this
phenomenon.
 The water molecules form a “cage”
around the small hydrophobic droplets.
Nonpolar Molecules
 Nonpolar molecules have no polar bonds or
the bond dipoles cancel due to molecular
geometry.
 These molecules do not form good
attractions with the water molecule. They
are insoluble and are said to be hydrophobic
(water hating).
 eg.: CH3CH2CH2CH2CH2CH3, hexane
Nonpolar Molecules-2
 Water forms hydrogen-bonded cage like structures
around hydrophobic molecules, forcing them out of
solution.
Amphipathic Molecules
 Amphipathic molecules contain both polar and
nonpolar groups.
 Ionized fatty acids are amphipathic. The
carboxylate group is water soluble and the long
carbon chain is not.
 Amphipathic molecules tend to form micelles,
colloidal aggregates with the charged “head”
facing outward to the water and the nonpolar “tail”
part inside.
A Micelle

Typical “Bond” Strengths
Type
kJ/mol
Covalent
>210
Non-covalent
Ionic interactions
4-80
Hydrogen bonds
12-30
van der Waals
0.3-9
Hydrophobic interactions
3-12
Osmosis, Reverse Osmosis & Dialysis
 Osmosis is a spontaneous process in which solvent
(e.g. water) molecules pass through a semi permeable
membrane from a solution of lower solute (e.g.
chemical) concentration to a solution of higher solute
concentration.
 Osmosis is the movement of solvent from a region of
high concentration (here, pure water) to a region of
relatively low concentration (water containing dissolved
solute).
 Water moves by osmosis and solutes by diffusion.
OSMOTIC PRESSURE
 Osmotic pressure is the pressure required to
stop osmosis or the influx of water (22.4 atm for
1M solution).
 Because cells have a higher ion concentration
than the surrounding fluids, they tend to pick up
water through the semi permeable cell
membrane.
 The cell is said to be hypertonic relative to the
surrounding fluid and will burst (hemolyze) if
osmotic control is not effected.
Cells placed in a hypotonic
solution will lose water and
shrink (crenate).
If cells are placed in an isotonic
solution (conc. same on both
sides of membrane) there is no
net passage of water.
Definitions of solutions
 Hypotonic solution: A solution with a lower
salt concentration than in normal cells of the
body and the blood.
 Hypertonic solution: A solution with a higher
salt concentration than in normal cells of the
body and the blood.
 Isotonic solution: A solution that has the same
salt concentration as the normal cells of the
body and the blood. An isotonic beverage may
be drunk to replace the fluid and minerals which
the body uses during physical activity.
OSMOMETER
 Osmotic pressure (p) is measured in an
osmometer.
Osmotic-pressure formula
p = iMRT
i = van’t Hoff factor (% as ions)
M = molarity (mol/L for dilute solutions)
R = (normal gas constant expressed in liters and
atmospheres)
0.082 L atm/ mol K
T = Kelvin temperature
Or
π= i*C*R*T
T= absolute temperature (in Kelvin)
R= the gas constant in whatever units you need to
express osmotic pressure (e.g. if you want π in atm
then R=0.082 L*atm/(mole*K))
C = the concentration of your solute in mole/L
 i is the van't Hoff coefficient. For non-
electrolytes i=1
 For strong electrolytes i= the number of ions
that are produced by the dissociation according
to the molecular formula
e.g for NaCl you have 2 ions (1 Na+ and 1 Cl-)
so i=2,
for CaCl2, 3 ions (1 Ca+2 from 2 Cl-) so i=3.
 For weak electrolytes, if n is the number of ions
coming from the 100% dissociation according to
the molecular formula and a the degree of
dissociation then i=(1-a)+na.
E.g. if we assume for CH3COOH a=80% i=(10.8)+2*0.8= 0.2+1.6=1.8
Liquids move from high osmotic pressure
(high conc. solvent and low conc. solute) to
low osmotic pressure (high conc. solute
and low conc. of solvent)
REVERSE OSMOSIS
 Reverse osmosis (RO) is a separation process
that uses pressure to force a solution through a
membrane that retains the solute on one side
and allows the pure solvent to pass to the other
side.
 More formally, it is the process of forcing a
solvent from a region of high solute
concentration through a membrane to a region
of low solute concentration by applying a
pressure in excess of the osmotic pressure.
 It is used in water purification and desalination.
DIALYSIS
 A concentrated solution is separated from a
large volume of solvent by a dialysis membrane
or bag that is permeable to both water and
solutes.
 Only small molecules can diffuse through the
pores of the membrane.
 At equilibrium, the concentrations of small
molecules are nearly the same on either side of
the membrane, whereas the macromolecules,
such as proteins or nucleic acids, remain inside
the dialysis bag.
KIDNEY DIALYSIS
 Reverse osmosis is the technique used
in dialysis, which is used by people with
kidney failure.
 The kidneys filter the blood, removing
waste products (e.g. urea) and water,
which is then excreted as urine.
 A dialysis machine mimics the function
of the kidneys. The blood passes from
the body via a catheter to the dialysis
machine, across an osmotic filter.
Sect 12: Water ionization, pH, titration and buffer
 The self-ionization of water is the chemical reaction in
which two water molecules react to produce a hydronium
(H3O+) and a hydroxide ion (OH−).
 Water ionization occurs endothermically due to electric
field fluctuations between molecules caused by nearby
dipole librations resulting from thermal effects, and
favorable localized hydrogen bonding.
 Ions may separate but normally recombine within a few
min. to seconds. Rarely (about once every eleven hours
per molecule at 25°C, or less than once a week at 0°C)
the localized hydrogen bonding arrangement breaks
before allowing the separated ions to return, and the pair
of ions (H+, OH-) hydrate independently and continue
their separate existence.
Ionization of Water
Water dissociates. (self-ionizes)
H2O + H2O =
H3O+ + OH-
Ka = [H3O+][OH-]
[H2O]2
Kw = Ka [H2O]2 = [H3O+ ][OH-]
Water Ionization-2
The conditions for the water
dissociation equilibrium must hold
under all situations at 25o C.
+
-14
Kw= [H3O ][OH ]=1 x 10
In neutral water,
+
-7
[H3O ] = [OH ] = 1 x 10 M
Water ionization - 3
When external acids or bases are
+
added to water, the ion product ([H3O
][OH ] ) must equal Kw.
The effect of added acids or bases is
best understood using the BronstedLowry- theory of acids and bases.
Bronsted-Lowry definitions
An acid is a substance that can donate a proton
A base is a substance that can accept a proton
H+ ions (called a protons, since a H+ ion has neither
electrons nor neutrons).
This definition can be represented by the general chemical
reaction
A  B + H+
which does not attempt to show electrical charge balance.
In this equation ·
A is the acid.,
·
B is the base and
·
H+ (a hydrogen atom without an electron) is a proton.
Conjugate acid/base
 An acid can donate a proton
 An acid (HA) reacts with a base (H2O) to
form the conjugate base of the acid (A-)
and the conjugate acid of base (H3O+)
HA + H2O = H3O+ + AA
B
CA
CB
C: conjugate (product) A/B
Conjugate base/acid
base = proton acceptor
RNH2 + H2O = OH- + RNH3+
B
A
CB
CA
Measuring Acidity
 Added acids increase the concentration of
hydronium ion and bases the concentration of
hydroxide ion.
 In acid solutions
[H3O+] > 1 x 10-7 M
[OH-] < 1 x 10-7 M
-7
 In basic solutions [OH ] > 1 x 10
M
+
-7
[H3O ] < 1 x 10 M
 pH scale measures acidity without using
exponential numbers.
pH Scale
+
Define: pH = - log(10)[H3O ]
0---------------7---------------14
acidic
basic
[H3O+]=1 x 10-7 M, pH = ?
7.0
Strength of Acids
Strength of an acid is measured by
the percent which reacts with water
to form hydronium ions.
Strong acids (and bases) ionize
close to 100%.
 e.g.. HCl, HBr, HNO3, H2SO4
Strength of Acids-2
Weak acids (or bases) ionize
typically in the 1-5% range .
e.g.. CH3COCOOH, pyruvic acid
CH3CHOHCOOH, lactic acid
CH3COOH, acetic acid
Strength of Acids-3
Strength of an acid is also measured
by its Ka or pKa values.
HA + H2O = H3O+ + AKa = [H3O+][A-]
[HA]
Larger Ka and smaller pKa values
indicate stronger acids.
Strength of Acids-4
Ka
pKa
CH3COCOOH
CH3CHOHCOOH
3.2x10-3
1.4x10-4
2.5
3.9
CH3COOH
1.8x10-5
4.8
Larger Ka and smaller pKa values indicate stronger acids.
Monitoring Acidity
The Henderson-Hasselbalch (HH)
equation is derived from the
equilibrium expression for a weak
acid.
[A ]
pH = pKa + log
[HA]
Monitoring Acidity-2
 The HH equation enables us to calculate the
pH during a titration and to make predictions
regarding buffer solutions.
 What is a titration?
It is a process in which carefully measured
volumes of a base are added to a solution of
an acid in order to determine the acid
concentration.
Monitoring Acidity-3
 When chemically equal (equivalent)
amounts of acid and base are present
during a titration, the equivalence point is
reached.
 The equivalence point is detected by using
an indicator chemical that changes color or
by following the pH of the reaction versus
added base, ie. a titration curve.
Glutamic acid titration curve
Titration Curve – solved example
 0.7 equivalents of NaOH neutralizes 0.7 eq
of acid producing 0.7 eq of salt and leaving
0.3 eq of unneutralized acid.
 pKa of HOAc is 4.76
pH = 4.76 + log [0.7]
[0.3]
30% acid and 70% salt.
pH=5.13
Buffer Solutions
 Buffer : a solution that resists change in pH
when small amounts of strong acid or base
are added.
 A buffer consists of:
a weak acid and its conjugate base or
 a weak base and its conjugate acid

Buffer Solutions - 1
High concentrations of acid and
conjugate base give a high
buffering capacity.
Buffer systems are chosen to
match the pH of the physiological
situation, usually around pH 7.
Buffer Solutions- 2
Within cells the primary buffer is the
phosphate buffer: H2PO4-/HPO42The primary blood buffer is the
bicarbonate system: HCO3-/H2CO3.
Proteins also provide buffer capacity.
Side chains can accept or donate
protons.
Buffer Solutions-3
A Zwitter ion is a compound with both
positive and negative charges.
Zwitterionic buffers have become
common because they are less likely to
cause complications with biochemical
reactions.
Example of Zwitter ion buffer
 N-tris(hydroxymethyl)methyl-2-
aminoethane sulfonate (TES) is a
example of zwitterion buffer.
+
(HOCH2)3CN H2CH2CH2SO3
Buffer Solutions-4
 Buffers work by chemically tying up acid and
base. Eg.:
HCO3- + H3O+
H2CO3 +
OH
H2CO3 + H2O
HCO3
+ H2O
Solved example on Buffer
Calculate the ratio of lactic acid to lactate in a buffer at
pH 5.00. The pKa for lactic acid is 3.86
5.00 = 3.86 + log [lactate]
[lactic acid]
5.00-3.86 = log [lactate]
[lactic acid]
antilog 1.14 = [lactate]
[lactic acid]
= 13.8
Sect 13: SUMMARY
 Water is essential for all living things.
 Water molecules can form hydrogen
bonds with other molecules because
they have 2 H atoms that can be
donated ans 2 unshared electron pairs
that can act as acceptors.
 Liquid water is an irregular network of
water molecules that each form 4
hydrogen bonds with neighbouring water
molecules.
Summary contd.
 Hydophilic substances such as ions and
polar molecules dissolve readily in water.
 The hydrophobic effect is the tendency
of water to minimizeits contact with
nonpolar substances.
 Water molecules move from regions of
high concentration to regions of low
concentration by osmosis.
 Solutes move from regions of high conc.
to regions of low conc. by diffusion.
Summary contd.
 Water ionizes to H+ (which represents the
hydronium ion H3O) and OH-.
 The concentration of H+ ions in solutions is
expressed as a pH value.
 Acids can donate protons and bases accept
protons.
 The strength of an acis is expressed as its pK.
 Henderson-Hasselbalch equation relates the
pH of a solution to the pK and concentration of
an acid to its conjugate base.
 Buffered solutions resist changes in pH within
about one pH unit of the pK of the buffering
species.
THE END Note
& Lao Tzu’s quote on water
 “Be careful what you water your
dreams with. Water them with worry
and fear and you will produce weeds
that choke the life from your dream.
Water them with optimism and
solutions and you will cultivate
success. Always be on the lookout for
ways to turn a problem into an
opportunity for success. Always be
on the lookout for ways to nurture
your dream.”
Tutorial on Water
1. Noncovalent bonding of water has play a
vital role in determining the properties of
water. Describe types of those bonding and
water properties they determined.
2. Differentiate polar and nonpolar
molecule.Explain how polar nature of water
makes it an excellent solvent for polar and
ionic materials.